Revision on the Atmoshper topic, Chemistry AS OCR B (Salters)

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Molecules and Networks

Gas Molecules move as far as part as possible- because they have weak intermolecular forces, only instantaneous dipole-induced dipole forces(ID- ID), therefore little energy energy is needed to separate them.

Molecules- have a specific formula

e.g carbon dioxide and oxygen.

Networks- no specific formula

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Networks- Silicon Oxide

Covalent Network e.g. SiO2, found as quartz or sand, empirical formula Properties:

  • Hard
  • High MP
  • Non-conductor of electricity- all of bonding its bonding are used for making covalent bonds.
  • Insoluble
  • Crystalline- regular repeating structure

As there covalent bonds in all directions. 


  • Tetrahedral
  • Silicon covalently bonded with four oxygen atoms. Oxygen is only bonded to two silicon.
  • Crystal lattice (Giant lattice)
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Molecules and Networks (2)

Molecules have:

  • Strong covalent bonds within molecules
  • Weak intermolecular forces between molecules
  • They require small amounts of energy for melting or boiling because ID- ID forces  have to be overcome. 
  • Sometimes permanent dipole forces operate e.g. hydrogen bonds in water

Networks have:

  • Strong covalent bonds between all atoms
  • They require large amounts of energy for melting as covalent bonds have to be broken. 
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Networks- Diamonds

Bonding- Covalently bonded to four other carbon atoms 

Shape- Tetrahedral shape, forms a crystal lattice shape

Because of its strong covalent bonds: 

  • Hard
  • Good thermal conductor, vibrations travel easily through the stiff lattice
  • High melting point- sublimes
  • Can't conduct electricity- all outer electrons are held in localised bonds
  • It won't dissolve in any solvent

Particularly hard- No weaknesses- bond lengths, strength, and bond numbers are all the same

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Graphite and Fullerenes

Graphite Network

  • Bonding:
    •  3 Covalent bonds to C atoms, 
    • 4th electron is delocalised between layers
  • Properties: 
    • conducts electricity and heat because of the delocalised electron, 
    • hard but slippery, 
    • high mp

Fullerenes e.g Buckmisterfullerene "Buckyball"

  • Molecule: C60
  • However can be made into a a network
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Why is bonding different in CO2 and SiO2?

Same group but act very differently in some ways. Similar in the fact that pure silicon has a similar structure to diamond. 

CO2- The O atom is held close to the small C atom's nucleus. Therefore, 2 oxygen electrons can be attracted and held in a double bond. 

"SiO2"-The Si atom is much bigger than a C atom, so its nucleus is further away from the O electrons. Only one electron per O atom can be attracted by each Si nucleus. There are 4 single bonds per Si atom

Covalent bonds between all atoms in Silicon (IV) oxide means that it has a very high mp. Whereas, carbon dioxide only has weak ID-ID  forces between molecules, therefore little energy is needed to separate its molecules 

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Reaction Rates and Temperature

A reaction won't take place between two particles unless-

  • They collide in the right direction. Facing each other the right way.
  • They collide with a least a certain minimum amount of kinetic energy- activation enthalpy.

Increasing the temperature:

  • More Kinetic Energy
  • A greater percentage of particles...
  • ... having at least the activation enthalpy
  • Therefore greater chance of successful collisions

Small increase in temperature, large increase in rate. 

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Rate of reaction and use of catalyst


  • increase rate of reaction,
  •  without being used up,
  •  it lowers activation enthalpy
  •  provide an alternative reaction pathway  with a lower activation enthalpy. 

If the activation enthalpy is lower more particles have enough enthalpy to react. 

Homogeneous catalysts- catalyst and reactants in same state e.g. enzymes in body cells, all aqueous.

Homogeneous catalysts speed up reactions by forming intermediate compounds. Products are formed from intermediate compounds. Therefore if homogeneous catalyst is used it will have two humps

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Collision Theory- The other factors

Increasing Concentration- Particles are closer together, therefore collide more often. More collisions mean more chances to react.

Increasing Pressure- Particles are closer together, collide more often, more chance to react.

Increasing Surface Area- If one reactant is in a lump then most of the particles won't collide with other reactants.  Crushing these lumps so that more of the particles can come in contact with other reactants. A smaller particle size means a larger surface area. This leads to a speedier reaction.   

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Dynamic Equilibrium

  • Goes both ways at the same speed.
  • Can only happen in a closed circuit

E.g. H2 (g)+ I2(g)-> 2HI(g)

Le Chatelier's Principle- tells you the position of equilibrium if conditions are changed

If change is imposed on a system- the equilibrium shifts to oppose change.

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Dynamic Equilibrium- Changing Pressure, Temperatur


The position of equilibrium will shift to side with fewer particles (if pressure is increased, reducing pressure)

Only affects equilibria involving gases.


  • Increasing, adding heat, equilibrium shifts in the endothermic (positive) direction to absorb this heat. 
  • Decreasing, removing heat, equilibrium shifts in the exothermic (negative) direction to try and replace heat

If forward is exothermic, backwards is endothermic.

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Dynamic Equilibrium- Changing Concentration and Us


  • Increase concentration reactants- equilibrium tries to get rid of the extra reactant by making more product. Equilibrium shifts to the right.
  • Increase the concentration of product- equilibrium tries to use up extra product. The reverse reaction goes faster, equilibrium shifts to to the left, decreasing concentration in opposite direction. 


Catalysts don't affect the position of equilibrium. Catalysts have no effect on the position of equilibrium-> CAN'T increase yield, but equilibrium is reached faster.

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The Atmosphere

Most of the atmosphere is Nitrogen and Oxygen.

  • Sun gives out electromagnetic radiation because of the nuclear processes going on its core.
    • Which is transmitted as waves, with a spectrum of differently  frequencies.
  • Sun mainly gives out visible(light), infra-red (heat) along with a smaller amount of UV radiation.
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The Atmosphere (2)

Energy is quantised, (exchanged in specific quantities-have certain vibrational energy levels)



Only frequencies of radiation corresponding to particular amount of energy are absorbed. Different molecules absorb different frequencies of radiation.

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The Atmosphere (3)

  • Sunscreens & ozone absorb some solar radiation
  • Light of the right frequency can break covalent bonds

UV- Ionisation

Visible- Dissociation

Infra-red- Rotation

Microwaves- Rotation

Radio waves- Translation

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Infra-red Radiation Makes Some Bonds Vibrate More

  • Some molecules absorb energy from infra-red radiation. The extra energy makes their covalent bonds vibrate more
  • Only molecules made of different atoms can absorb infra-red radiation, carbon dioxide, water, nitrogen oxide and methane do. Oxygen and nitrogen don't. Gases that  absorb infra-red are called greenhouse gases. 
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UV and Visible Light Give Electrons more energy

The electrons in molecules are quantised.

  • When UV or Visible Light hits a molecule or gas the electrons can absorb the energy and jump up to their next energy level, because the energy needed for these changes is quantised too, only specific frequencies are absorbed. 
  • If enough energy is absorbed bonds break, forming free radicals
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Radical Reactions


Heterolytic fission- Different ions, two different substances are formed- a cation and a negatively charged cation.

Homolytic fission- Radicals formed

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Radical Reactions (2)

Initiation- something that produces more radicals than is used up (when radicals are produced)

Propagation- Creates as many radicals as they use up.

Terminations- the end.

  • Radicals are reactive, particles with unpaired electron.
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Greenhouse Effect

  • Radiation from the Sun reaches the Earth's atmosphere, most of the UV and infra-red is absorbed by atmospheric gases and some radiation is reflected back into space from clouds.
  • The energy that reaches the Earth's surface is visible light, mainly & UV. Some is reflected into space by light-coloured, shiny surfaces like ice & snow. The rest is absorbed by the Earth, which causes it to warm up. 
  • The earth then radiates energy back towards space as infra-red radiation (heat). -> Some escapes (through the so called 'IR window'- range where frequency are not absorbed by atmospheric gases') and re-emit it in all directions- including towards Earth, keeping us warm.
    • Various gases in the troposphere absorb other infra-red radiation and re-emit it in all directions- including back towards earth keeping us warm.- Greenhouse effect
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Contribution of a particular gas depends on:

  • How much radiation one molecule of the gas absorbs
  • How much of that gas there is in the atmosphere

Limit Global Warming: 

  • Use renewable fuels e.g. wind, solar, hydrolic
  • Artificial photosynthesis proteins
  • Increase Cost
  • Population control
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Halogenalkanes and CFCs

  • Halogens react with alkanes into photochemical reactions. Photochemical reactions are started by ultraviolet light.
  • A hydrogen atom is substituted by chlorine and bromine. A free-radical substitution reaction.

Initiation reactions- sunlight provides enough to break the Cl-Cl - this photo-dissociation. The bond splits equally and each atom gets to keep one electron-homolytic fission. The atom becomes a highly reactive free radical, because of its unpaired electron

Propagation reactions- free radicals are used up and created in a chain reaction. Chlorine radical attacks a methane molecule. 

The new methyl free radical, can attack another chlorine molecule

The new chlorine molecule can attack a methane molecule, and so on, until all the chlorine or methane molecules are wiped out.

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Halogenoalkanes and CFCs- Termination Reactions

The free radicals are mopped up.

  • If two free radicals join together, they make a stable molecule.
  • There are heaps of possible termination reactions. 
    • Examples:
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UV Radiation Can Also Break Carbon-Halogen Bonds

  • All halogenoalkanes contain bonds between carbon atoms and halogen atoms. Ultraviolet radiation can break these bonds- the carbon-halogen bond splits homolytically to create two free radicals
  • The ease with which the carbon-halogen bond is broken by UV depends on the halogen. It turns out that the carbon-iodine bond is the most likely to break, and the carbon-fluorine bond the least likely. This is because the C-I bond has the lowest bond enthalpy, and C-F bond the highest.
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Halogenoalkanes and CFCs (3)

  • Chlorofluorocarbons(CFCs) are a group of compounds made by replacing all of the hydrogen atoms in alkanes with chlorine and fluorine. They're halogenoalkanes.
  • They're unreactive, non-flammable and harmless. It's also possible to make CFCs with a range of boiling points. Used in Fire extinguishers, aerosols, coolant in fridges. 

But CFCs help create 'holes' in the ozone layer of the Earth's atmosphere.

Ozone holes still form in spring but the rate of decrease of ozone is slowing.

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Ozone is in the stratosphere-> Ozone is formed when UV radiation from the Sun hits the oxygen molecules.

Ozone should be destroyed be destroyed and replaced as UV radiation hits molecules. Equilibrium is set up:

However formation of ozone is slow.

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Ozone (2)

Ozone layer which is is high in atmosphere protects against UV radiation.

  • Ozone layer removes all the high energy, frequency UVC radiation ,and 90% of UVB, which are harmful.

Scientists found 'holes' in Ozone, which is bad as it lets more UV radiation reach earth.

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CFCs destroying the ozone layer

CFCs absorb energy in the stratosphere producing Cl (homolytic fission)

Free radicals act as catalysts.

Overall reaction                                                    

and Cl as catalyst depletion of  ozone.

Nitrogen oxides destroy ozone in the same way as radicals.

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Ozone in the Troposphere

Ozone occurs in the troposphere, due to the effect of sunlight on mixtures of nitrogen dioxide and hydrocarbons. These occur naturally from a variety of source but vehicle engines and powerstations contribute a large amount too.

-> Ozone in the troposphere is toxic.

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